Final answer:
Valence bond theory accounts for the trend of decreasing bond angles in CH4, NH3, and H2O by considering hybridization and electron pair repulsion, which explain the deviations from idealized predictions based on simple atomic orbital overlaps.
Step-by-step explanation:
Valence Bond Theory and Bond Angles
The question asks how valence bond theory accounts for the trend of decreasing bond angles in CH4 (methane), NH3 (ammonia), and H2O (water). The answer lies in the concept of hybridization of atomic orbitals. Methane, with a tetrahedral geometry, has bond angles of 109.5°, as expected from the sp3 hybridization of carbon's valence orbitals. In ammonia, one of these hybrid orbitals contains a lone pair of electrons, resulting in a slight distortion to 107.5° due to the lone pair repulsion. In water, two of the sp3 orbitals hold lone pairs, further reducing the bond angle to 104.5°. These observations are due to variations in the way atomic orbitals combine and the repulsion between electron pairs. Without considering hybridization and electron pair repulsion, the predictions of bond angles from the simple overlap of atomic orbitals would not match the observed values.
Furthermore, the size of the central atom affects the degree of orbital overlap. In the case of water, the combination of oxygen's 2s and 2p orbitals results in hybrid orbitals that are less than the ideal 109.5° due to lone pair shaping of the molecular geometry. This contrasts with the predictions of simple 2p orbital overlap which fails to adequately describe the real structure.